A STUDY  OF  THE  DETERMINATION  OF  BROMIDES 
IN  THE  PRESENCE  OF  EXCESS  CHLORIDES 

BY 

MAYOR  FARTHING  FOGLER 
B.  S.  University  of  Illinois,  1920 


THESIS 

Submitted  in  Partial  Fulfillment  of  the  Requirements  for  the 

Degree  of 

MASTER  OF  SCIENCE 
IN  CHEMISTRY 

IN 

THE  GRADUATE  SCHOOL 
OF  THE 

UNIVERSITY  OF  ILLINOIS 


1921 


. 


■ 


UNIVERSITY  OF  ILLINOIS 


THE  GRADUATE  SCHOOL 


August  1. 192-1  • 


I HEREBY  RECOMMEND  THAT  THE  THESIS  PREPARED  UNDER  MY 

SUPERVISION  BY -Mayor  Farthing  Fogler 

ENTITLED A Study  of  the  Deterrri nation  of  Bromides  in 

the  Presence  of  Excess  Chlorides 

BE  ACCEPTED  AS  FULFILLING  THIS  PART  OF  THE  REQUIREMENTS  FOR 
THE  DEGREE  OF has  tan  of  Science^ 


Recommendation  concurred  in* 


Committee 

on 

Final  Examination* 


^Required  for  doctor’s  degree  hut  not  for  master’s 


476970 


A STUDY  OF  THE  DETERMINATION  OF  BROMIDES 
IN  THE  PRESENCE  OF  EXCESS  CHLORIDES 


MAYOR  FARTHING  FOGLER 
B.  S.  University  of  Illinois,  1920 


THESIS 

Submitted  in  Partial  Fulfillment  of  the  Requirements  for  the 

Degree  of 

MASTER  OF  SCIENCE 
IN  CHEMISTRY 

IN 

THE  GRADUATE  SCHOOL 
OF  THE 

UNIVERSITY  OF  ILLINOIS 


1921 


' Vi*  to  rr  ffc  rmr4  ;r  • >iti  rcfutt 


If  <>TAIJC1AM{>  till 

"tl  .O 

ic  hi  >f  '-i  * n (Vi nu 


Table  of  Contents. 


Page 

I . Introduction . ...  1 


II.  Historical  .......... 

Ill  . Experimental 

A.  Gravimetric  Methods  ...... 

Precipitation  of  Silver  Bromide  in  the 
Presence  of  Excess  of  Chlorides 
Precipitation  from  Ammonium  Hydroxide 
Solutions  ....... 

Alcohol  as  a Solvent  * 

Cause  of  Low  Results  ..... 

Empirical  Corrections  for  Low  Results 


2 

4 


4 


5 

7 

S 

9 


B.  Selective  Oxidation  Methods  . 

Sodium  Sulfite  as  a Reducing  Agent 
Oxidizing  Agents  ...... 

Fixation  of  Bromine  by  Metallic  Silver 
Theory  of  Oxidation  of  Bromides  and 

Chlorides  

Method  Recommended  ..... 


14 

15 

16 
19 

19 

21 


IV.  Summary ....  23 

V.  Bibliography  25 


Plates . 


1. 

2. 


Opposite 

Page 


Solubility  of  AgBr  in  NaCl  Solution  ...  13 

Change  of  Oxidation  Potential  with  Concentration  20 


Digitized  by  the  Internet  Archive 
in  2016 


https://archive.org/details/studyofdeterminaOOfogl 


Acknowledgment 


The  writer  wishes  to  express 
his  appreciation  and  sincere 
thanks  to  Doctor  J.  H.  Reedy 
for  his  assistance  in  the 
experimental  work,  and  also 
for  his  aid  in  writing  this 
thesi s . 


A Study  of  the  Determination  of  Bromides 
In  the  Presence  of  Excess 
of  Chlorides . 

I.  Introduction* 


A fundamental  problem,  in  the  bromine  industry  is  the 

accurate  estimation  of  the  bromide  content  of  the  brines  or 
* 

bitterns.  Since  the  bromide  content  is  low  usually  less 
than  5^  — this  problem  involves  the  determination  of  bro- 
mides in  the  presence  of  a large  excess  of  chlorides.  It 
is  important  because  it  is  desirable  to  know  just  how  much 
bromine  is  present,  so  as  to  add  exactly  the  equivalent 
amount  of  oxidizing  agent  needed  to  effect  its  liberation. 
Various  methods  in  current  use  which  give  satisfactory  re- 
sults for  high  concentrations  give  very  divergent  results 
for  low  concentrations.  The  purpose  of  this  investigation 
is  to  study  the  methods  for  estimating  bromine  and  to  ascer- 
tain, if  possible,  the  cause  of  their  inaccuracy. 

II  . Historical ♦ 

Chemical  literature  contains  numerous  methods  that  have 
been  proposed  for  this  analysis,  most  of  them  depending  on 


(1) 


. 

© 

. 


■ 


p 

a selective  oxidation  of  the  bromide  and  its  subsequent  titra- 
tion or  gravimetric  determination.  Vortman3-  who  attempted 
to  determine  chlorine  in  the  presence  of  bromine  used  lead 
dioxide  in  the  presence  of  acetic  acid  to  oxidise  the  bro- 
mides and  then  determined  gravimetrically  the  chlorine  which 
remained  in  the  solution.  The  results  were  always  high  be- 
cause the  bromine  was  not  all  set  free,  while  if  a higher 
concentration  of  acetic  acid  was  used  traces  of  chlorine 
were  liberated.  Cavazzis  used  barium  peroxide  and  sulfuric 
acid  as  the  oxidizing  agent.  The  bromine  was  distilled,  but 
a little  chlorine  was  also  liberated.  Eng.le^  tried  to  use 

ammonium  persulfate  and  sodium  acetate  at  70°  -80°.  Berg- 
4 

lund  used  potassium  hydrogen  sulfate  and  potassium  perman- 
ganate in  the  cold,  and  aspirated  air  through  the  solution 
to  remove  the  bromine,  collecting  it  in  sodium  hydroxide 
solution.  He  reports  that  the  bromine  was  then  determined 
gravimetrically,  though  no  statement  is  made  as  to  how  the 
sodium  hypobromite  formed  was  reduced.  It  was  found,  more- 
over, that  while  the  oxidizing  agent  would  not  liberate  chlo- 
rine if  it  alone  were  present,  it  did  liberate  chlorine  when 
bromine  was  present.  He  avoided  this  by  a double  aspiration. 
Baubigny  and  Rivals0  used  copper  sulfate  and  potassium  per- 
manganate at  the  constant  temperature  of  £0°,  but  with  little 
success,  as  chlorine  was  also  liberated.  Wyss^  attempted 
to  separate  all  of  the  halogens  by  selective  oxidation.  He 
first  removed  the  iodine  by  ferric  sulfate,  and  then  added 


. 

' 

. 

- 

■ 


. 


3 


potassium  permanganate  and  Seated  to  60°  to  liberate  the  bro- 
mine which  was  determined  gravimetrically . Ee  stated  that 
his  method  gave  good  results,  but  offered  no  figures  to  sup- 
port his  claim.  White^  used  aluminium  sulfate  and  potassium 
permanganate  and  stated  that  bromine  was  liberated  while  chlo 
rine  and  iodine  were  not.  It  is  quite  inconceivable  how  an 
oxidizing  agent  could  have  sufficient  oxidizing  power  to  lib- 
erate bromine  and  at  the  same  time  not  liberate  iodine  which 
is  much  more  readily  oxidized.  Jannasch  and  Aschoff-  used 
acetic  acid  and  permanganate  tc  liberate  the  bromine  but 
their  results  were  always  io w.  Bugarszky^  very  carefully 
investigated  other  methods  proposed,  and  himself  used  iodic 
acid  to  free  bromine,  but  obtained  very  poor  results.  An- 
drews1^ used  iodic  acid  to  oxidize  the  bromide  and  found  it 
very  suitable  if  the  bromine  was  present  in  large  quantities. 
But  this  method  is  hardly  applicable  where  chlorine  is  pres- 
ent in  large  quantities.  Later  investigators  have  found 
out  that  the  oxidation  potential  of  the  oxidizing  agent  must 
lie  between  that  of  bromine  and  chlorine.  As  an  oxidizing 
agent  of  the  correct  potential  iodic  acid  has  been  used,  as 
mentioned  above.  Gooch11  used  selenic  and  telluric  acids, 
but  neither  works  well  unless  chlorine  is  present  in  very 
minute  amounts.  Skinner  and  Baughman155  used  hydrogen  per- 
oxide and  chromic  acid  to  liberate  the  bromine,  which  was 
removed  by  aspiration  in  the  cold.  They  found  that  chlorine 
also  came  over  from  a saturated  solution  to  the  extent  of 


. 


' 

. 

- 


4 


about  ifo.  This  was  remedied  by  double  aspiration.  They  ob- 
tained very  good  results,  but  the  method  is  rather  long  and 
involved. 

None  of  the  above  methods  are  adapted  to  rapid,  accu- 
rate work,  and  seme  of  them  admit  of  errors  as  large  as  2*fc 
Those  that  are  fairly  accurate  require  several  hours  or  even 
days  for  their  completion.  It  would  appear,  then,  that  the 
optimum  method  has  not  yet  been  developed. 

III.  Experimental . 

There  are  in  general  two  methods  for  determining  bro- 
mides in  the  presence  of  an  excess  of  chlorides:  (1)  The 

gravimetric  method,  which  involves  the  precipitation  of  the 
bromine  and  all  or  part  of  the  chlorine  as  silver  halides 
and  the  subsequent  estimation  of  the  composition  of  the  pre- 
cipitate, either  by  indirect  analysis  or  calculation;  ( 2 ) 

The  selective  oxidation  of  the  bromine,  using  an  oxidizing 
agent  that  will  not  liberate  chlorine,  or  at  least  only  a 
small  amount. 

A.  Gravimetric  Methods . 

Precipitation  of  Silver  Bromide  in  the  Presence  of  Ex- 
cess of  Chlorldes--A  gravimetric  method  that  appears  — ■ on 
its  fac6,  at  least  --  to  be  both  rapid  and  accurate,  involves 
the  addition  of  a known  amount  of  silver  nitrate  to  the  ha- 
lide solution  in  a quantity  somewhat  in  excess  of  that  needed 
to  completely  precipitate  the  bromine.  The  amount  of  bro- 


- 


. 

. 


. 

. 

— 


5 


mine  present  may  be  calculated  from  th©  following  expression: 

Wt.  of  Br  = 1.7976  x Wt . of  AgBr  + AgCl  - 

2.3885  x wt.  of  Ag. 


The  excess  of  silver  nitrate  over  the  amount  necessary  to 
preceipitate  the  bromine  should  be  small.  This  follows  from 
the  fact  that,  for  mixtures  of  silver  bromide  and  silver 
Chloride, 


$ AgBr  = 100-=  - 13..00  x-  ..-?!■£ 


Wt.  of  AgBr  + AgCl 


Now,  in  case  of  an  error  in  the  weight  of  the  silver  halide 
precipitate,  the  effect  cn  the  result  will  be  small  if  the 
amount  of  silver  chloride  is  small;  but  if  the  latter  is 
large,  the  value  of  the  fraction  in  the  above  expression 
will  be  increased  considerably;  that  is,  the  magnitude  of 
the  inaccuracy  will  be  multiplied  as  many  fold,  as  the  weight 
of  the  silver  chloride  is  increased. 

The  accuracy  of  this  method  depends  on  the  assumption 
that  the  bromine  is  quantitatively  precipitated  by  the  sil- 
ver nitrate  before  appreciable  amounts  of  silver  chloride 
are  formed.  This  inference  in  justified  by  a comparison  of 
the  solubility  products  of  the  two  halides. 

Precipitation  from  Ammonium  Hydroxide  Solutions--! t 
was  felt  that  the  above  method  was  open  to  criticism  in  that, 
owing  to  the  speed  of  the  reaction,  the  silver  nitrate  might 
be  used  up  locally  in  the  formation  of  silver  chloride,  leav- 
ing unprecipitated  bromides  in  the  solution.  It  is  known 


- 


. 


: 


- 


- 

- 


6 


that  the  conversion  of  silver  chloride  into  silver  bromide 
by  means  of  soluble  bromides  is  low.  Hence  if  the  precipi- 
tation could  be  made  to  proceed  slowly,  the  silver  bromide 
precipitation  would  occur  first,  and  the  silver  chloride  pre- 
cipitation would  not  begin  until  the  former  is  complete. 

In  order  to  do  this  it  is  necessary  to  use  a solvent  for  the 
silver  halides  that  may  be  gradually  removed,  or  else  to 
add  the  silver  nitrate  very  slowly.  Ammonium  hydroxide  was 
the  first  solvent  to  be  used.  The  experiments  were  carried 
out  as  follows:  Solutions  of  potassium  bromide  and  potaseium 

chloride  were  mixed,  a volume  of  ammonium  hydroxide  added, 
and  then  enough  silver  nitrate  to  precipitate  all  the  bro- 
mine and  part  of  the  chlorine.  The  beaker  containing  the 
mixture  was  heated  to  boiling  and  stirred  vigorously  until 
all  the  ammonia  was  driven  out.  Any  traces  were  neutralized 
by  the  addition  of  dilute  nitric  acid,  and  the  precipitate 
filtered  in  a Gooch  crucible,  dried  and  weighed.  Table  I 
shows  typical  results. 


Table  I_.  Results  with  Ammonium  Hydroxide  as  Solvent . 


Exp. 

Per  cent 
KOI 

Amount  of  AgBr 
Theory 

Amount  of  AgBr 
Found 

Difference 

1 

Sfo 

1.5480  gm. 

1.5254  gm. 

-.0126  gm 

2 ! 

2$ 

1 . 5480 

1.4850 

-.0630 

5 

2 f0 

1.5480 

1.5405 

-.0075 

4 

2 fo 

1.5480 

1.4945 

-.0535 

5 

P% 

1 . 5480 

1.4861 

-.0619 

6 

P$ 

1 . 5480 

1.5533 

-.0147 

7 

1 . 5480 

1.4666 

-.0814 

8 

9$ 

1 . 5480 

1.4260 

-.1220 

9 

1.5480 

1.4789 

-.0691 

10 

9$ 

1.5480 

1.4996 

-.0484 

11 

2% 

1.5480 

1.4771 

-.0709 

* * I 1C1  Bkff-  ;’|H^ 

' ' 


. 

. 

. 

— 

— 

« . 

> ■— 

o . 

. 

• X 

• - 

. 

■ 

. 

. 

7 


It  will  be  noted  that  the  results  were  always  low,  indi- 
cating incomplete  precipitation  of  either  the  bromine  or  the 
silver.  This  fact  will  be  interpreted  later. 


Alcohol  as  a Solvent — It  was  thought  that  the  silver 
halides  were  both  tcc  soluble  in  ammonium  hydroxide  to  get 
consistent  results,  so  some  other  solvent  was  sought  for. 


It  was  observed  that  silver  chloride  is  appreciably  soluble 


in  ethyl  alcohol,  while  silver  bromide  is  not.  So  the  above 
determinations  were  repeated,  except  50$  alcohol  was  used 
as  the  solvent.  Table  II  sho?/s  the  result?.. 

Table  II.  Results  with  50$  Alcohol  as  Solvent. 


Exp . 

Per  cent 
KC1 

Amount  of  AgBr 
Theory 

Amount  of  AgBr 
Found 

Difference 

1 

2$ 

1.5480  gm. 

1.5518  gm. 

+.0038 

2 

2$ 

1.5480 

1.5158 

-.0312 

3 

2$ 

1.5480 

1.4956 

-.0524 

4 

2$ 

1.5480 

1.5231 

-.0249 

5 

2$ 

1.5480 

1.5080 

-.0400 

6 

2$ 

1.5480 

1.5100 

-.0380 

7 

2% 

1.5480 

1.5268 

-.0212 

8 

2%. 

1.5480 

1.5328 

-.0152 

9 

2$ 

1 . 5480 

1.5121 

-.0359 

10 

2$ 

1.5480 

1.5334 

-.0146 

Water  as  a 

Solvent- -Mainly 

for  the  purpose 

i 

•H 

u 

cfi 

P* 

a 

o 

o 

o 

son 

with  the  above  results,  a series  of  determinations  was 

made 

with  water 

as  the  solvent. 

Table  III  shows 

the  results. 

Table  III.  Results  with  Water  as  Solvent. 

Exp. 

Per  cent 
KC1 

Amount  of  AgEr 
Theory 

Amount  of  AgBr 
Found 

Difference 

1 

2$ 

1.5480 

1.5228 

-.0252 

s 

2% 

1.5480 

1.5250 

-.0230 

3 

2$ 

1.5480 

1.5344 

-.0126 

4 

2$ 

1.5480 

1.5098 

-.0382 

. 


. 

• 

. 

_ 

. 

_ 

• — 

. 

. 

♦ _ .-j 


• 

- 

. 

, 

The  results  in  the  experiments  in  water  solution  are 
distinctly  better  than  with  ammonium  hydroxide  and  alcohol 
as  solvents,  but  still  they  fall  far  below  the  usual  stan- 
dard for  accuracy. 

To  meet  the  suggestion  that  the  low  results  were  due 
to  the  bromine  not  being  completely  precipitated,  owing  to 
the  slow  conversion  of  silver  chloride  into  silver  bromide, 
the  following  experiment  was  made:  The  silver  in  4 oc.  of 

normal  silver  nitrate  was  completely  precipitated  as  sil- 
ver chloride  and  then  stirred  with  an  excess  of  sodium  bro- 
mide solution  for  six  hours,  after  which  it  was  found  that 
about  75 °jo  of  the  chlorine  had  been  replaced  by  bromine.  In 
the  precipitation  of  silver  from  silver  nitrate  solution 
by  mixed  halides,  on  the  other  hand,  the  silver  halide  is 
formed  in  a very  highly  dispersed  condition,  which  has  a 
far  higher  solubility  than  the  coagulated  form,  and  in  this 
state  the  bromine  seems  to  replace  the  chlorine  in  a com- 
paratively  short  time.  It  would  appear  that  the  magnitude 
of  these  dispersed  particles  is  so  small  that  no  particle 
is  large  enough  to  permit  of  the  formation  of  a protective 
coating  of  silver  bromide  and  thus  prevent  the  inner  part 
from  reacting  with  the  sodium  bromide. 

Cause  of  Low  Re suit a- -The  cause  of  the  low  results  in 
the  precipiation  methods  is  to  be  found  in  the  solvent  action 
of  alxaline  halides  on  silver  bromide  and  silver  chloride. 
Schierhclz-1-0  has  reported  that  100  grams  of  sodium  chloride 
in  concentrated  solution  dissolve  0.474  grams  of  silver  bro- 


- 


. ) 


- 


9 


mide,  and  IOC  grams  of  potassium  bromide  in  concentrated  solu- 
tion dissolve  3.019  grams  of  this  salt,  the  temperature  being 
15°  in  both  cases.  In  a similar  way,  silver  chloride  is 
also  dissolved  by  solutions  of  concentrated  alkaline  halides. 


Empirical  Corrections  for  Low  Result s--The  expedient 
of  deriving  a set  of  solubility  values  for  silver  bromide 
in  alkaline  halides  was  then  taken  up.  Given  these  data, 
empirical  corrections  could  be  made  on  all  analytical  re- 
sults. So  it  was  planned  to  construct  a solubility  curve 
for  silver  bromide  in  concentrations  of  sodium  chloride  vary- 
ing from  if*  to  saturation.  The  determinations  were  made  as 
fellows:  50  cc . of  approximately  0.1  N silver  nitrate  were 

taken  and  just  enough  sodium  bromide  was  added  to  precipi- 
tate all  of  the  silver  as  silver  bromide.  Results  are  given 
in  Table  IV. 

Table  IV . Solubili ty  of  Silver  Bromide  in  Sodium 

Chloride  Solution. 


Grams  per 
100  cc » of 

Weight  of  AgBr 

Weight  of  AgBr 

Weight  of  AgBr 

Solution 

Theory 

Found 

Dissolved 

1 

.9354 

.9178 

.0176 

1 

.9354 

.9100 

.0254 

1 

.9354 

.9169 

.0185 

1 

.9324 

.9204 

.0120 

1 

.9324 

.9193 

.0131 

1 

.9324 

.9200 

.0124 

1 

.9798 

.9706 

.0092 

1 

.9798 

.9705 

.0093 

1 

.9798 

.9712 

.0086 

2 

.9354 

.9087 

.0267 

2 

.9354 

.9044 

.0310 

2 

.9354 

.9087 

.0267 

-• 

r ' 

* 


. 

. 


. 


. 


Table  IV--Cont. 


Grams  per 
100  cc « cf 

Weight  of  AgBr 

Weight  of  AgBr 

Weight  of  AgB 

-L-  V V w v1  • V*  i* 

Solution 

Theory 

Found 

Dissolved 

2 

.9324 

.9133 

.0191 

2 

.9324 

.9195 

.0129 

2 

.9324 

.9143 

.0181 

2 

.9798 

.9668 

.0130 

2 

.9798 

.9683 

.0115 

2 

.9798 

.9656 

.0142 

3 

.9354 

.9012 

.0342 

3 

.9354 

.90  §5 

.0329 

3 

.9354 

.9027 

.0327 

3 

.9324 

.9082 

.0242 

3 

.9324 

.9100 

.0226 

3 

.9324 

.9077 

.0247 

4 

.9354 

.8986 

.0368 

4 

.9354 

.8959 

.0395 

4 

.9354 

.8976 

.0378 

4 

.9324 

.9055 

.0269 

4 

.9324 

.9060 

.0264 

4 

.9798 

.9605 

.0293 

4 

.9798 

.9597 

.0301 

4 

.9798 

.9588 

.0316 

5 

.9354 

.8900 

.0454 

5 

.9354 

.8908 

.0446 

5 

.9354 

.8905 

.0449 

5 

.9324 

.9014 

.0310 

5 

.9324 

.9022 

.0302 

6 

.9354 

.8905 

.0449 

6 

.9354 

.8902 

.0446 

6 

.9354 

.8900 

.0448 

6 

.9324 

.9000 

.0324 

6 

.9324 

.8994 

.0330 

7 

.9354 

.8871 

.0483 

7 

.9354 

.8885 

.0469 

7 

.9324 

.8958 

.0368 

7 

.9324 

.8958 

.0368 

7 

.9324 

.8952 

.0374 

. 


. 


. 


. 


11 


Table  IV-=»Ccnt. 


Grains  per 
1GG  cc.  of 
Solution 

Weight  of  AgBr 

Weight  of  AgBr 

Weight  of  AgBr 

Theory 

Found 

Dissolved 

7 

.9798 

.9537 

.0867 

7 

.9798 

.9580 

.0878 

7 

.9798 

.9586 

.0878 

8 

.9354 

.8830 

.0511 

8 

.9354 

.8836 

.0505 

8 

.9384 

.8945 

.0379 

8 

.9384 

.8945 

.0379 

8 

.9384 

.8939 

.0385 

9 

.9354 

.8888 

.0538 

9 

.9354 

.8815 

.0539 

9 

.9354 

.8817 

.0537 

9 

.9798 

.9473 

.0385 

9 

.9798 

.9507 

.0891 

9 

.9798 

.9458 

.0340 

10 

.9354 

.8771 

.0581 

10 

.9354 

.8790 

.0568 

10 

.9354 

.8784 

.0568 

10 

.9758 

.9444 

.Q354 

10 

.9758 

.9444 

.0354 

10 

.9758 

.9440 

.0358 

11 

.9798 

.9400 

.0398 

11 

.9798 

.9416 

.0388 

11 

.9798 

.9389 

.0409 

18 

.9798 

.9396 

.0408 

IS 

.9798 

.9389 

.0409 

17 

.9798 

.9888 

.0570 

17 

.9798 

.9888 

.0516 

17 

.9798 

.9890 

.0508 

SO 

.9798 

.9185 

.0613 

SO 

.9798 

.9184 

.0614 

SO 

.9798 

,9141 

.0657 

ss 

.9798 

.9140 

.0656 

ss 

.9798 

.9187 

.0671 

ss 

.9798 

.9189 

.0669 

12 


Table  IV— Cent. 


Grams  per 
ICC1  cc . of 
Solution 

Weight  of  AgBr  Weight  of  AgEr 
Theory  Found 

Weight  of  AgBr 
Dissolved 

24 

.9798 

. .9079 

.0719 

24 

.9798 

.9127 

.0671 

24 

.9798 

.9120 

.0678 

26 

.9798 

.9042 

.0756 

26 

.9798 

.9040 

.0758 

26 

.9798 

.908  2 

.0716 

28 

.9798 

.8924 

.0874 

28 

.9798 

.8909 

.0887 

28 

.9798 

.8937 

.0861 

30 

.9798 

.8830 

.0968 

30 

.9798 

.8805 

.0993 

30 

.9798 

.8797 

.1001 

32 

.9798 

.8681 

.1117 

32 

.9798 

.8699 

.1099 

34 

.9798 

.8540 

.1258 

34 

.9798 

.8591 

.1207 

34 

.9798 

.8538 

,1260 

36 

.9798 

.8408 

.1390 

36 

.9798 

.8409 

.1389 

36 

.9798 

.8450 

.1348 

In  the 

above  table 

the  results  from 

one  and  the  same 

solution  are 

set  off  in 

blocks 

of  three, 

or 

sometimes  two. 

It  will  be  noticed  that, 

for  the  most  part, 

there  is  good 

agreement  within  these  blocks. 

For  different  blocks,  how- 

ever,  wrhere 

solutions  of 

sodium 

bromide  i 

of 

different  con- 

centrations 

were  used  in 

preparing  the  silver  bromide,  and 

where  uniform  conditions 

were  not  maintained,  the  results 

are  markedly 

different . 

Take , 

for  example. 

the  solubilities 

in  1 % sodium 

chloride  solution. 

The  first 

three  detemiina- 

tions  are  from  the  same 

concentration  of 

sodium  bromide,  were 

run  at  the  same  time,  and  under  conditions  that  were  iden 


30 

36 

34 

32 

30 

26 

26 

24 

22 

20 

/8 

/ 6 

J4 

/2 

/ 0 

6 

6 

4 

2 

0 


G/7)<s.  /VaC/  per  /OO  cc  /VaO 


— ] 

13 

tical,  including  concentration,  heating,  stirring,  and  so 
forth . 

This  indefiniteness  in  the  solubility  of  silver  bromide 
is  explained  by  the  fact  that  substance  is  a colloid,  and 
the  size  of  the  particles  depends  on  such  conditions  as  those 
just  mentioned.  And  in  turn,  the  solubility  of  a solid  de- 
pends on  the  size  of  the  particles,  the  solubility  increas- 
ing enormously  with  the  degree  of  subdivision. 

Attempts  were  made  to  regulate  the  conditions  by  carry- 
ing out  all  determinations  alike.  For  example,  each  precip- 
itate was  stirred  for  a definite  time  to  coagulate  the  par- 
ticles, and  washed  with  a definite  amount  of  water.  In  this 
way  better  checks  were  obtained.  If  a solution  of  sodium 
bromide  of  the  same  concentration  we re  used  throughout,  good 
results  could  probably  be  obtained  and  a fairly  smooth  solu- 
bility curve  drawn  for  the  solubility  of  silver  bromide  in 
various  strengths  of  sodium,  chloride  solution.  Figure  I 
shows  three  solubility  curves  plotted  from  the  data  in  Table 
IV.  Curve  I was  plotted  from  solubility  measurements  on 
silver  bromide  made  by  mixing  .0994  N sodium  bromide  and 
.1043  N silver  nitrate  in  equivalent  amounts;  curve  II  from 
.1007  N sodium  bromide  and  .0993  N silver  nitrate;  curve 
III  from  .1006  N sodium  bromide  and  .0996  N silver  nitrate. 

Other  conditions  than  concentration  were  practically  the 
same  in  all  cases. 

It  will  be  seen  therefore  that  the  amount  of  silver 
bromide  which  dissolves  in  a certain  concentration  of  sodium 


' 


. 


- 

- 


. 


, • I 1 1 


14 


chloride  solution  is  not  constant  unless  all  details  of  tech- 
nique are  regulated  so  as  to  be  exactly  identical  in  every 
case.  These,  of  course,  can  not  be  regulated  on  an  unknown 
sample.  Therefore  it  is  not  possible  to  determine  bromine 
in  a solution  of  this  kind  by  precipitating  and  weighing  it 
and  then  adding  an  empirical  correction  which  corresponds 
to  its  solubility  in  that  concentration  of  sodium  chloride 
solution . 

£ • Selective  Oxidation  Methods . 

Methods  involving  the  selective  oxidation  of  bromides 
to  free  bromine  and  the  subsequent  distillation  of  the  lib- 
erated halogen  offer  certain  advantages,  the  most  important 
of  which  is  the  concentration  of  the  bromine.  Later  work 
has  shown  that  it  is  not  always  possible  to  make  a "clean” 
separation  from  chlorine  in  this  way,  since  varying  amounts 
of  the  latter  are  usually  present  in  the  distillate.  But 
it  does  have  the  advantage  of  reducing  the  chloride  concen- 
tration very  considerably.  The  importance  of  this  point 
is  evident  when  it  is  remembered  that  the  halogen  content 
of  the  distillate  is  almost  always  determined  gravimetrically , 
and,  as  has  been  shown  above,  the  gravimetric  method  has  the 
maximum  accuracy  when  the  amount  of  chlorine  is  a minimum. 

The  whole  range  of  oxidizing  agents  has  been  explored 
by  investigators  in  their  efforts  to  find  one  that  will  pref- 
erentially oxidize  the  bromine  without  affecting  the  chlo- 
rine. Nitrosyl  sulfuric  acid  ("ni trose" ) , chromic  acid, 
acidified  potassium  permanganate,  and  so  forth,  have  been 


. 


. 


15 


used  with  more  cr  less  uncertain  success. 

An  important  point  in  the  technique  of  these  distilla- 
tion methods  is  that  the  halogens  of  the  distillate  must 
he  absorbed  in  some  suitable  reducing  agent  which  will  re- 
convert them  into  bromides  and  chlorides,  respectively.  If 
silver  nitrate  were  also  present  in  this  absorbing  medium, 
the  halogens  would  be  immediately  precipitated  in  their 
final  form  for  weighing.  The  use  of  silver  nitrate  without 
a reducing  agent  would  lead  to  low  results,  owing  to  the 
formation  of  oxidized  bromine  compounds  like  hypobromous 
acid  and  silver  bromate. 


Sodium  Sulfite  as  a Reducing  Agent- -The  first  agent 
tried  was  sodium  sulfite.  The  determination  was  carried 
out  as  follows:  5G  cc . of  0.1  N sodium  bromide  was  placed 

in  a distilling  flask  and  an  amount  of  sodium  chloride  was 
added  and  the  solution  diluted  to  about  100  cc.  To  this 
was  added  IS  grams  of  chromic  acid  and  about  2 cc.  of  30$ 
hydrogen  peroxide,  and  the  mixture  distilled.  The  solution 
in  which  the  halogens  were  to  be  absorbed  consisted  of  5 
cc.  of  0.1  N silver  nitrate  and  the  calculated  amount  of 
sodium  sulfite.  The  distillation  was  run  for  about  30  min- 
utes. The  precipitate  was  coagulated  by  stirring,  collected 
in  a Gooch  crucible,  dried  and  weighed.  Results  were  as 
follows: 


Br  Present 


Br  Found 


1.0562  gm 
1.0562 


9999  gm 
9853 


- 


f 


■ 


. 


■ 


16 


The  results  in  these  determinations  were  low  because  the 
sodium  sulfite  reduced  some  of  the  silver  nitrate  to  free 
silver  on  heating.  In  order  to  avoid  this  complication  a 
series  of  experiments  were  conducted  in  which  the  absorbing 
solution  was  not  heated.  Results  were: 

Br  Present  Br  Found 


.9865  gm 
.9865 
.9865 
.9865 


.9930  gm, 
.9870 
.9943 
.9898 


These  values  are  high  --  presumably  because  varying  amounts 
of  chlorine  were  carried  over  during  the  distillation. 

The  next  determinations  were  made  by  absorbing  the 
bromine  in  sodium  sulfite  solution,  then  acidifying  with 
sulfuric  acid  and  boiling,  so  as  to  expell  the  sulfur  diox- 
ide. Silver  nitrate  solution  was  then  added,  and  the  bro- 
mine determined  in  the  usual  way.  Only  a small  excess  of 
oxidizing  agent  was  used  in  the  reaction  flask  beyond  what 
was  theoretically  required  to  liberate  the  bromine,  so  as 
not  to  oxidize  the  chlorine  if  possible.  Results: 

Br  Present  Br  Found 

.9865  gto.  .9822  gm. 

.9865  .9843 

These  determinations  show  that  the  oxidation  of  the  bromides 
was  not  complete,  and  that  an  excess  of  oxidizing  agent  is 
necessary  to  effect  its  quantitative  expulsion. 

Cxi dizing  Agents — The  chromic  acid-hydrogen  peroxide 
mixture  is  not  a very  satisfactory  oxidizing  medium,  be- 
cause if  the  mixture  is  heated  long  enough  to  drive  off  all 


. 

. 


. 

. 


. 


. 

. 


- 

. • 


17 

of  the  bromine  large  amounts  of  chromyl  chloride  distill 
over,  and  thus  too  large  an  excess  of  silver  nitrate  must 
be  used  to  precipitate  all  of  the  halogen  present. 

Another  oxidizing  agent  experimented  with  was  hydrogen 
peroxide  in  the  presence  of  sulfuric  acid.  It  was  observed 
that  hydrogen  peroxide  in  the  presence  of  concentrated  sul- 
furic acid  will  liberate  bromine  from  sodium  bromide  solu- 
tion at  50°  - 50°,  or  even  a lower  temperature,  while  the 
chlorine  in  sodium  chloride  solution  appears  unaffected  until 
the  solution  is  brought  to  a boil.  This  would  suggest  that 
it  might  be  possible  to  liberate  bromine  almost  quantitatively 
from  a solution  of  the  halides  by  this  reagent.  In  order  to 
ascertain  if  all  the  bromide  could  be  oxidized  quickly  and 
also  to  find  out  if  it  would  be  completed  absorbed  in  sil- 
ver nitrate  solution  to  which  hydrogen  peroxide  had  been 
added  to  act  as  a reducing  agent,  a series  of  experiments 
was  made  in  which  no  chlorine  at  all  was  present.  The  reac- 
tions involved  are  as  follows: 

S JSaBr  + Hg02  + H2 S04  -*  Na2SG4  + 2 H20  + Br2. 

This  is  an  oxidation  reaction,  the  hydrogen  peroxide  acting 
as  the  oxidizing  agent.  In  the  following  reaction  it  con- 
ducts itself  as  a reducing  agent: 

P.  AgNOg  + Erg  + HgGg  -»  P AgBr  + 2 BNOg  + 0g. 

These  experiments  were  carried  out  in  the  following  manner: 

20  cc.  of  the  sodium  bromide  solution  was  placed  in  an  as- 
pirator bottle  and  2 cc.  of  30 ^ hydrogen  peroxide  and  5 cc. 
of  concentrated  sulfuric  acid  added.  The  liberated  bromine 


- 


18 


was  then  removed  by  aspiration  and  collected  in  nitrogen  bulbs 
containing  25  cc.  of  dilute  silver  nitrate  solution  and  about 
2 cc.  of  30$  hydrogen  peroxide.  The  solution  in  the  aspira- 
tor bottle  became  colorless  after  aspiration  for  about  one- 
half  hour.  The  results  were  as  follows: 


Br  Present 

Br  Pound 

.4012  gm. 
.4012 
.4012 
.4012 

.4027  gm. 
.4013 
.4018 
.4023 

Why  these  results  should  be 

somewhat  high  has  not  been  ac- 

counted  for.  However,  it  should  be  noted  that  this  method 
gives  decidedly  the  most  satisfactory  results  of  all  those 
tried. 

The  following  determinations  were  made  on  solutions  of 
sodium  bromide  in  7$  sodium  chloride,  using  only  a slight  ex 
ce3s  of  silver  nitrate  to  collect  the  bromine: 


Br  Present 

Br  Found 

.4012  gm. 

.4012 

.4012 

.3656  gm. 

.3683 

.3678 

The  following  were  made  using  a larger  excess  of  silver  ni 
trate: 


Br  Present 

Br  Pound 

,4012  gm. 
.4012 
.4012 
.4012 

.4119  gm. 
.4147 
.4110 
.4163 

This  demonstrates  the  fact  that  even  in  fairly  dilute  solu- 
tions of  sodium  chloride  the  sulfuric  acid  will  liberate 
hydrogen  chloride,  which  may  be  more  or  less  oxidized  to 


• 

. 

' 

. '1  - 

. 

- ..  . 


.f 

. 

. 


. 

. 


19 


free  chlorine.  At  any  rate  the  silver  precipitate  contains 
silver  chloride. 

Fixation  of  Bromine  b£  Metallic  Sjiver--An  interesting 
possibility  suggested  itself  in  the  expedient  of  absorbing  the 
bromine  by  passing  the  aspirated  gases  through  a tube  con- 
taining finely  divided  silver.  The  principle  upon  which  this 
conjecture  was  based  is  the  very  general  belief  that  bromine 
will  react  with  metallic  silver  while  hydrochloric  acid  will 
not.  The  procedure  was  as  follows:  Very  fine  crystals  of 

pure  silver  were  prepared  by  electrolyzing  silver  nitrate 
between  a silver  anode  and  a platinum  cathode.  These  were 
carefully  washed  and  dried,  and  placed  in  a tube  through 
which  the  gases  were  to  be  drawn.  Upon  trial,  however,  tak- 
ing no  precautions  to  dry  the  gases,  the  increase  in  weight 
was  too  high,  indicating  that  moist  hydrochloric  acid  will 
react  with  fine  divided  silver.  Upon  introducing  between 
the  tube  and  the  flask  wash  bottles  containing  concentrated 
sulfuric  acid  to  dry  the  gases,  the  increase  in  weight  was 
practically  nil,  showing  that  silver  will  not  combine  with 
dry  bromine. 

Theory  of  Oxidation  of  Bromides  and  Chlorides- -The  fact 
that  no  one  has  found  an  oxidizing  agent  that  will  selectively 
oxidize  the  bromine  finds  a complete  explanation  in  thermo- 
dynamic considerations.  According  to  thermodynamics,  the 
chemical  work  done  in  these  oxidations  depends  wholly  upon 
the  "initial  and  final  states"  and  is  "independent  of  the 


Loy<y/-/  f/?/77s  of  O/  ft/ f/o/?s 

/oo  /ooo  / 


A60 


20 

path"  --  that  is,  the  oxidizing  agent  used.  By  "initial  and 
final  states"  is  meant  particularly  concentration  conditions* 
This  doctrine  teaches,  therefore,  that  the  expenditure  of 
energy  in  effecting  the  oxidation  of  a certain  amount  of 
bromide  ions  at  definite  concentration  to  bromine  is  a con- 
stant, and  that  a certain  "thermodynamic  potential"  is  nec- 
essary for  the  process.  As  a measure  of  this  thermodynamic 
potential  may  be  used  the  electrical  potential  necessary  to 
bring  about  electrically  the  oxidation,  which  of  course  is 
equal  to  the  change  in  "free  energy"  involved.  It  is  pos- 
sible, of  course,  that  a particular  oxidizing  agent  might  be 
somewhat  "faster"  than  another  agent,  and  in  this  way  show 
selective  action.  But  for  similar  ions  like  Br“  and  01”, 
this  is  not  likely. 

Figure  II  shows  the  relation  between  the  concentrations 
of  the  halogen  ions,  Br"  and  Cl”,  and  the  oxidizing  poten- 
tial necessary  to  convert  them  into  free  halogens.  These 
graphs  have  been  calculated  by  means  of  the  well-known 
Nernst  formula  from  the  values  of  the  normal  bromine  and 
chloine  electrodes.  The  temperature  is  assumed  to  be  °5°  C. 

The  horizontal  axis  shows  the  dilutions  (that  is,  the  recip- 
rocals of  the  concentrations)  expressed  in  terms  of  their 
logarithms.  Using  0.1  milligram  of  bromine  as  the  limit  of 
analytical  error,  the  potential  necessary  to  reduce  the  Br”- 
ion  concentration  (using  100  cc . of  solution)  to  that  point 
has  been  calculated  by  the  figure  to  be  1.36  volts,  referred 
to  the  normal  hydrogen  electrode  as  0.0  volts.  This  potential, 


.• 


- 


'■ 


■ 


* 


21 

however,  as  also  shown  by  the  figure,  is  sufficient  to  oxi- 
dize Cl“=ions  of  concentrations  greater  than  0.6  N sodium 
chloride  --  that  is,  a concentration  of  3.5 <f>  by  weight. 

This  shows  very  decidedly  why  selective  oxidation  methods 
will  not  bring  about  a complete  separation  of  bromine  from 
moderate  concentrations  of  sodium  chloride  like  5 $ --  which 
is  a weak  strength  as  brines  and  bitterns  go.  For  higher  con' 
centrations,  approaching  saturated  sodium  chloride  solution, 
liberation  of  the  chlorine  may  begin  when  100  cc.  of  the 
solution  still  contain  as  much  as  4.5  milligrams  of  unoxi- 
dized bromide. 

Method  Recommended--The  results  obtained  in  thi3  work 
as  discussed  above  indicate  very  conclusively  that  bromine 
can  not  be  separated  quantitatively  from  an  excess  of  chlo- 
rides by  selective  oxidation  and  distillation,  since  oxida- 
tion of  the  chloride  begins  before  the  bromine  is  completely 
expelled.  The  best  method  is  the  following:  The  halide  mix- 

ture is  placed  in  an  aspirator  bottle  and  2 cc.  of  30^  hy- 
drogen peroxide  and  about  5 cc.  of  concentrated  sulfuric 
acid  added.  The  halogens  are  collected  in  a known  amount 
of  dilute  silver  nitrate  containing  a volume  of  pure  hydro- 
gen peroxide  --  for  example,  2,  cc.  of  the  3 Oft  product.  After 
the  aspiration  has  continued  for  about  three  quarters  of  an 
hour,  the  vessel  containing  the  absorbing  medium  is  discon- 
nected and  placed  on  a water  bath  until  the  solution  "bright- 
ens." Now  dilute  potassium  chloride  solution  is  carefully 
added  drop  by  drop  as  long  as  a precipitate  forms.  The  pre- 


~ 


- 

■ 


22, 

cipitat©  is  then  treated  in  the  usual  way,  and  its  composi- 
tion calculated  by  the  expression  on  page  5. 


23 


IV . Summary, 

I,  Indirect  methods  for  calculating  bromides  in  the 
presence  of  chlorides  are  accurate  only  when  the  chloride 
concentration  is  low. 

P.  Slow  precipitation  of  silver  halides  from  ammonium 
hydroxide  and  from  alcoholic  solutions  give  unsatisfactory 
results  in  the  presence  of  excess  chlorides. 

3.  Low  results  for  bromine  are  due  to  the  solubility 
of  silver  bromide  in  alkaline  halides. 

4.  Empirical  corrections  for  low  results  are  not  feas- 
ible, since  the  solubility  of  silver  bromide  varies  with 
its  state  of  subdivision. 

5.  In  selective  oxidation  processes  the  halogens  are 
distilled  into  an  absorption  medium;  sodium  sulfite  was  stu- 
died as  a reducing  agent  for  this  medium  and  proved  trouble- 
some by  reducing  the  silver  nitrate  at  high  temperatures. 

6.  Hydrogen  peroxide  was  found  very  satisfactory  as 
an  oxidizing  agent  for  liberating  the  bromine,  and  as  a 
reducing  agent  for  reconverting  it  to  the  bromide  form. 

7.  An  excess  of  oxidizing  agent  is  desirable  in  the 
liberation  of  the  bromine  from  the  halide  solution. 

8.  There  is  always  a simultaneous  liberation  of  chlo- 
rine along  with  the  bromine  in  the  oxidation  of  solutions  of 
bromides  and  chlorides,  particularly  if  the  latter  are  pres- 
ent in  quantity. 

9.  The  simultaneous  liberation  of  bromine  and  chlorine 


. 


1 


24 


is  explained  from  thermodynamic  considerations,  which  indi- 
cate that  there  is  no  oxidizing  agent  that  will  effect  a 
satisfactory  selective  oxidation  of  bromine  in  the  presence 
of  chlorides  in  excess. 

10.  A new  procedure  of  estimating  bromine  in  the  pre3 
ence  of  chlorides  is  suggested. 


25 


V.  Bibliography. 

1.  Vortman:  Z.  anal.  Chem.,  25,  172  (1886). 

2.  Cavazzi:  Gazz.  Chim.  Ital.,  13,  174. 

3.  Rngle:  Compt.  Rend.,  118,  1263  (1894). 

4.  Berglund:  Z.  anal.  Cham,,  24,  184  (1885). 

5.  Baubigny  and  Rivals:  Compt.  Rend.,  125,  527  (1827). 

5.  Wyes:  Rept.  anal.  Chem.,  5,  238  (1885). 

7.  White:  Chem.  News,  1,  14 4 (1892). 

8.  Jannasch  and  Aschoff:  Z.  anorg.  Chem.,  1,  144  (1892). 

9.  'Bugarszky:  Z.  anorg.  Chem.,  10,  387  (1897). 

10.  Andrews:  J.  Am.  C.  S.,  25,  809  (1903). 

11.  Gooch:  J.  Am.  C.  S.,  29,  275  (1907);  Am.  J.  Sci.,  35, 

54  (1913). 

12.  Skinner  and  Baughman:  J,  Ind.  Sng.  Chem.,  11,  954  (1919). 

13.  Schierholz:  Sitzber.  K.  Akad.  Wi3s.  (Vienna),  101,  2b, 

4 (1890). 

14.  Abegg,  Auerbach  and  Luther:  Messungen  elektromotori scher 

Kraefte  galvani scher  Ketten  (1911),  p.  200. 


